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Chapter 12 Electrons in Atoms

Chapter 12 Electrons in Atoms. Brodersen Honors Chem 2013/14. Bohr’s Model. Why don’t the electrons fall into the nucleus? Move like planets around the sun. In circular orbits at different levels. Amounts of energy separate one level from another. Bohr’s Model. Nucleus. Nucleus.

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Chapter 12 Electrons in Atoms

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  1. Chapter 12 Electrons in Atoms Brodersen Honors Chem 2013/14

  2. Bohr’s Model • Why don’t the electrons fall into the nucleus? • Move like planets around the sun. • In circular orbits at different levels. • Amounts of energy separate one level from another.

  3. Bohr’s Model Nucleus Nucleus Electron Electron Orbit Orbit Energy Levels Energy Levels

  4. Bohr postulated that: • Fixed energy related to the orbit • Electrons cannot exist between orbits • The higher the energy level, the further it is away from the nucleus • An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive)

  5. ELECTROMAGNETIC RADIATION

  6. Electromagnetic radiation.

  7. Electromagnetic Radiation Most subatomic particles behave as PARTICLES and obey the physics of waves.

  8. wavelength Visible light Amplitude wavelength Node Ultaviolet radiation Electromagnetic Radiation

  9. Electromagnetic Radiation • Waves have a frequency • Use the Greek letter “nu”, , for frequency, and units are “cycles per sec” • All radiation:  •  = cwhere c = velocity of light = 3.00 x 108 m/sec

  10. How did he develop his theory? • He used mathematics to explain the visible spectrum of hydrogen gas • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf

  11. High energy Low energy Low Frequency High Frequency X-Rays Radiowaves Microwaves Ultra-violet GammaRays Infrared . Long Wavelength Short Wavelength Visible Light

  12. The line spectrum • electricity passed through a gaseous element emits light at a certain wavelength • Can be seen when passed through a prism • Every gas has a unique pattern (color)

  13. Line spectrum of various elements Helium Carbon

  14. Bohr’s Triumph • His theory helped to explain periodic law • Halogens are so reactive because it has one e- less than a full outer orbital • Alkali metals are also reactive because they have only one e- in outer orbital

  15. Drawback • Bohr’s theory did not explain or show the shape or the path traveled by the electrons. • His theory could only explain hydrogen and not the more complex atoms

  16. Atomic Line Emission Spectra and Niels Bohr Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of theLINE EMISSION SPECTRAof excited atoms. • Problem is that the model only works for H Niels Bohr (1885-1962)

  17. Atomic Spectra One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit.

  18. Atomic Spectra and Bohr Bohr said classical view is wrong. Need a new theory — now called QUANTUM or WAVE MECHANICS. e- can only exist in certain discrete orbits e- is restricted to QUANTIZED energy state (quanta = bundles of energy)

  19. Quantum or Wave Mechanics Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms. He developed the WAVE EQUATION Solution gives set of math expressions called WAVE FUNCTIONS,  Each describes an allowed energy state of an e- E. Schrodinger 1887-1961

  20. Heisenberg Uncertainty Principle Problem of defining nature of electrons in atoms solved by W. Heisenberg. Cannot simultaneously define the position and momentum (= m•v) of an electron. We define e- energy exactly but accept limitation that we do not know exact position. W. Heisenberg 1901-1976

  21. Arrangement of Electrons in Atoms Electrons in atoms are arranged as LEVELS (n) SUBLEVELS (l) ORBITALS (ml)

  22. QUANTUM NUMBERS The shape, size, and energy of each orbital is a function of 3 quantum numbers which describe the location of an electron within an atom or ion n (principal) ---> energy level l (orbital) ---> shape of orbital ml(magnetic) ---> designates a particular suborbital The fourth quantum number is not derived from the wave function s(spin) ---> spin of the electron (clockwise or counterclockwise: ½ or – ½)

  23. QUANTUM NUMBERS So… if two electrons are in the same place at the same time, they must be repelling, so at least the spin quantum number is different! The Pauli Exclusion Principle says that no two electrons within an atom (or ion) can have the same four quantum numbers. If two electrons are in the same energy level, the same sublevel, and the same orbital, they must repel. Think of the 4 quantum numbers as the address of an electron… Country > State > City > Street

  24. Energy Levels } • Further away from the nucleus means more energy. • There is no “in between” energy • Energy Levels Fifth Fourth Third Increasing energy Second First

  25. The Quantum Mechanical Model • Energy is quantized. It comes in chunks. • A quanta is the amount of energy needed to move from one energy level to another. • Since the energy of an atom is never “in between” there must be a quantum leap in energy. • Schrödinger derived an equation that described the energy and position of the electrons in an atom

  26. Atomic Orbitals • Principal Quantum Number (n) = the energy level of the electron. • Within each energy level the complex math of Schrödinger's equation describes several shapes. • These are called atomic orbitals • Regions where there is a high probability of finding an electron

  27. S orbitals • 1 s orbital for every energy level 1s 2s 3s • Spherical shaped • Each s orbital can hold 2 electrons • Called the 1s, 2s, 3s, etc.. orbitals

  28. P orbitals • Start at the second energy level • 3 different directions • 3 different shapes • Each orbital can hold 2 electrons

  29. The p Sublevel has 3 p orbitals

  30. The D sublevel contains 5 D orbitals • The D sublevel starts in the 3rd energy level • 5 different shapes (orbitals) • Each orbital can hold 2 electrons

  31. The F sublevel has 7 F orbitals • The F sublevel starts in the fourth energy level • The F sublevel has seven different shapes (orbitals) • 2 electrons per orbital

  32. Summary Starts at energy level

  33. Electron Configurations • The way electrons are arranged in atoms. • Aufbau principle- electrons enter the lowest energy first. • This causes difficulties because of the overlap of orbitals of different energies. • Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

  34. Electron Configurations First Energy Level • only s sublevel (1 s orbital) • only 2 electrons • 1s2 Second Energy Level • s and p sublevels (s and p orbitals are available) • 2 in s, 6 in p • 2s22p6 • 8 total electrons

  35. Third energy level • s, p, and d orbitals • 2 in s, 6 in p, and 10 in d • 3s23p63d10 • 18 total electrons Fourth energy level • s,p,d, and f orbitals • 2 in s, 6 in p, 10 in d, and 14 in f • 4s24p64d104f14 • 32 total electrons

  36. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s

  37. Electron Configuration • Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to .

  38. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The first to electrons go into the 1s orbital • Notice the opposite spins • only 13 more

  39. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2s orbital • only 11 more

  40. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2p orbital • only 5 more

  41. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 3s orbital • only 3 more

  42. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The last three electrons go into the 3p orbitals. • They each go into separate shapes • 3 unpaired electrons • 1s22s22p63s23p3

  43. Orbitals fill in order • Lowest energy to higher energy. • Adding electrons can change the energy of the orbital. • Half filled orbitals have a lower energy. • Makes them more stable. • Changes the filling order

  44. Write these electron configurations • Titanium - 22 electrons • 1s22s22p63s23p64s23d2 • Vanadium - 23 electrons 1s22s22p63s23p64s23d3 • Chromium - 24 electrons • 1s22s22p63s23p64s23d4 is expected • But this is wrong!!

  45. Chromium is actually • 1s22s22p63s23p64s13d5 • Why? • This gives us two half filled orbitals. • Slightly lower in energy. • The same principal applies to copper.

  46. Copper’s electron configuration • Copper has 29 electrons so we expect • 1s22s22p63s23p64s23d9 • But the actual configuration is • 1s22s22p63s23p64s13d10 • This gives one filled orbital and one half filled orbital. • Remember these exceptions

  47. Great site to practice and instantly see results for electron configuration.

  48. Practice • Time to practice on your own filling up electron configurations. • Do electron configurations for the first 20 elements on the periodic table.

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